Complex ions precipitates, solubility

Actinide(III) precipitates dissolve to a limited extent in solution containing 1M (or more) (NH J CO. However, the actinide (I I I) carbonato complex ion precipitates very slowly by adding hexamminecobalt(III) chloride and the yield of precipitation is not high. Separations of actinide(IV) and (VI) ions from actinide (III) and (V) ions are thus achieved by taking advantage of their different solubilities in ammonium carbonate solution. Hexamminecobalt (III) salt is used as a precipitant to recover U(VI), Pu(IV) and Am(VI) ions from ammonium carbonate solution. 

Another interesting effect of complex ions on solubilities is illustrated by the addition of iodide ion to a solution containing mercury(II) ion. After a moderate amount of iodide ion has been added, an orange precipitate forms (Fig. 16.15) through the reaction... 

Equilibria Involving Compbx Ions 643 Formation of Complex Ions 643 Complex Ions and Solubility of Precipitates 645 CHAPTER REVIEW GUIDE 646 PROBLEMS 648... 

Silver Thiosulfate. Silver thiosulfate [23149-52-2], Ag 2 y is an insoluble precipitate formed when a soluble thiosulfate reacts with an excess of silver nitrate. In order to minimize the formation of silver sulfide, the silver ion can be complexed by haUdes before the addition of the thiosulfate solution. In the presence of excess thiosulfate, the very soluble Ag2(S203) 3 and Ag2(S203) 3 complexes form. These soluble thiosulfate complexes, which are very stable, are the basis of photographic fixers. Silver thiosulfate complexes are oxidized to form silver sulfide, sulfate, and elemental sulfur (see Thiosulfates). 

When a metal ion is chelated by a ligand such as citric acid, it is no longer free to undergo many of its chemical reactions. A metal ion that is normally colored may, in the presence of citrate, have Httie or no color. Under pH conditions that may precipitate a metal hydroxide, the citrate complex may be soluble. Organic molecules that are catalyticaHy decomposed in the presence of metal ions can be made stable by chelating the metal ions with citric acid. 

Metathesis reactions are sometimes the reverse of those in aqueous systems because of the differing solubility relations. For example because AgBr forms the complex ion [Ag(NH3)2]" " in liquid NH3 it is readily soluble, whereas BaBr2 is not, and can be precipitated ... 

The precipitate redissolves in excess soluble cyanide, and the complex ion is probably an ion-dipole co-ordination compound, i.e. 

The increase in solubility of a precipitate upon the addition of excess of the precipitating agent is frequently due to the formation of a complex ion. A... 

This complex ion dissociates to give silver ions, since the addition of sulphide ions yields a precipitate of silver sulphide (solubility product 1.6 x 10 49 mol3 L 3), and also silver is deposited from the complex cyanide solution upon electrolysis. The complex ion thus dissociates in accordance with the equation ... 

At this point, the silver and mercury(I) chlorides remain as precipitates. When aqueous ammonia is added to the solid mixture, the silver precipitate dissolves as the soluble complex ion Ag(NH3)2+ forms ... 

Both Co3+ and Ni2+ form soluble complex ions with ammonia ligands rather than hydroxide precipitates. 

Polyion complex technique [40] is a unique method for immobilization of bilayer membranes with polymers. Water-insoluble complex is precipitated as the polyion complex when an aqueous solution of the charged bilayer membrane is mixed with a water solution of the counter charged polyelectrolyte. Stoichiometric ion pair formation is often found. Aging of the precipitate in a hot mixture kept above phase transition temperature of the bilayer membrane completes the ion exchange reaction [41], Chloroform solution of the polyion complex is washed by water several times to remove water soluble components [42]. 

Copper forms practically aU its stable compounds in -i-l and +2 valence states. The metal oxidizes readily to -i-l state in the presence of various com-plexing or precipitating reactants. However, in aqueous solutions +2 state is more stable than -i-l. Only in the presence of ammonia, cyanide ion, chloride ion, or some other complexing group in aqueous solution, is the +1 valence state (cuprous form) more stable then the +2 (cupric form). Water-soluble copper compounds are, therefore, mostly cupric unless complexing ions or molecules are present in the system. The conversion of cuprous to cupric state and metalhc copper in aqueous media (ionic reaction, 2Cu+ — Cu° -i- Cu2+) has a Kvalue of 1.2x106 at 25°C. 

A) The reaction that produced the precipitate is Pb (a,) + 2C1 (aq) PbCl2(,). Lead chloride is a slightly soluble salt, with a solubility of 10 g/L at 20°C. The solubility of PbCF increases very rapidly as the temperature rises. At 100°C it has a solubility of 33.5 g/L. However, PbCl2 precipitates very slowly, particularly when other ions that form insoluble chlorides are not present. PbCb dissolves in excess chloride ion as a result of the formation of a complex ion, tetrachloroplumbate(II) ion ... 

The fate of heavy metals in aquatic systems depends on partitioning between soluble and partieulate solid phases. Adsorption, precipitation, coprecipitation, and complexation are processes that affect partitioning. These same processes, which are influenced by pH, redox potential, the ionic strength of the water, the concentration of complexing ions, and the metal concentration and type, affect the adsorption of heavy metals to soil (Richter and Theis 1980). 

If anion X- precipitates metal M+, it is often observed that a high concentration of X causes solid MX to redissolve. The increased solubility arises from the formation of complex ions, such as MX2, which consist of two or more simple ions bonded to each other. 

The solubility product is the equilibrium constant for the dissolution of a solid salt into its constituent ions in aqueous solution. The common ion effect is the observation that, if one of the ions of that salt is already present in the solution, the solubility of a salt is decreased. Sometimes, we can selectively precipitate one ion from a solution containing other ions by adding a suitable counterion. At high concentration of ligand, a precipitated metal ion may redissolve by forming soluble complex ions. In a metal-ion complex, the metal is a Lewis acid (electron pair acceptor) and the ligand is a Lewis base (electron pair donor). 

If there is introduced into the solution from some other source an ion that is in common with an ion of the insoluble solid, the chemical equilibrium is shifted to the left, and the solubility of that solid will be greatly decreased from what it is in pure water. This is called the 11 common-ion effect." This effect is important in gravimetric analysis, where one wishes to precipitate essentially all of the ion being analyzed for, by adding an excess of the "common-ion" precipitating reagent. There is a practical limit to the excess, however, which involves such factors as purity of precipitate and possibility of complex formation. You can calculate the solubility under a variety of conditions, as illustrated in the following problem. 

When KI is added to a solution of palladium(II) chloride, an insoluble diiodide is precipitated. The dark red-black crystals are soluble in excess iodide with form a (ion of rhe tetraiodide complex ion. Palladinm(TT) iodide evolves iodine at 100°C, the decomposition to the elements being complete at 330-360°C, The black compound, palladium(III) fluoride, is made by direct combination of the elements. On rednetion. the brown difluoiide is formed,... 

Because, of the veiy heavy ionic, weight (250) of the perrhenate ion, it is one of the heaviest simple anions obtainable in readily soluble salts. It has found use as a precipitant for potassium and some other heavy univalent ions also as a precipitant for such complex ions as Co(NH )( "+. and for the separation of alkaloids and organic bases. Perrhenate also is used in the fractional crystallization ot the rare-earth elements. 

In fresh water, silver may form complex ions with chlorides, ammonium (in areas of maximum biological activity), and sulfates form soluble organic compounds such as the acetate and the tartrate become adsorbed onto humic complexes and suspended particulates and become incorporated into, or adsorbed onto, aquatic biota (Boyle 1968). Where decaying animal and plant material are abundant, silver strongly precipitates as the sulfide or combines with humic materials (Smith and Carson 1977). 

OH, together with non-ionised NH4OH and molecular NHg, the fact that zinc hydroxide, precipitated by addition of ammonium hydroxide to a solution of the chloride, is re-dissolved by further addition of ammonia, is doubtless to be explained by the formation of the complex ion Zn(NH8)2, which is soluble in water. But this does not... 

Subject areas for the Series include solutions of electrolytes, liquid mixtures, chemical equilibria in solution, acid-base equilibria, vapour-liquid equilibria, liquid-liquid equilibria, solid-liquid equilibria, equilibria in analytical chemistry, dissolution of gases in liquids, dissolution and precipitation, solubility in cryogenic solvents, molten salt systems, solubility measurement techniques, solid solutions, reactions within the solid phase, ion transport reactions away from the interface (i.e. in homogeneous, bulk systems), liquid crystalline systems, solutions of macrocyclic compounds (including macrocyclic electrolytes), polymer systems, molecular dynamic simulations, structural chemistry of liquids and solutions, predictive techniques for properties of solutions, complex and multi-component solutions applications, of solution chemistry to materials and metallurgy (oxide solutions, alloys, mattes etc.), medical aspects of solubility, and environmental issues involving solution phenomena and homogeneous component phenomena. 

Many models include features allowing calculations involving the speciation of soluble complexes, the precipitation and dissolution of solid mineral phases and the adsorption of ions from solution on to surfaces. In addition, conditions of oxidation-reduction and partial pressures of gas phases can be superimposed on the calculation. Throughout the chapter some simple examples of chemical modelling calculations will be given. 

Just as acids may be used to lower the concentration of anions in solution, so complexing agents may be used in some cases to lower the concentration of cations. In this problem, the addition of ammonia converts most of the silver to the complex ion, [Ag(NH3 )2]+. The upper limit for uncomplexed [Ag+] without formation of a precipitate can be calculated from the solubility product. 

Obviously one could measure the pH of a known concentration of a weak acid and obtain a value of its hydronium ion activity, which would permit a direct evaluation of its dissociation constant. However, this would be a one-point evaluation and subject to greater errors than by titrating the acid halfway to the equivalence point. The latter approach uses a well-buffered region where the pH measurement represents the average of a large number of data points. Similar arguments can be made for the evaluation of solubility products and stability constants of complex ions. The appropriate expression for the evaluation of solubility products again is based on the half-equivalence point of the titration curve for the particular precipitation reaction [AgI(OH2)2h represents the titrant] ... 

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